1. IB Topics 9 & 19 AP Chapters 4.9-4.10; 17
Oxidation & Reduction
IB Topics 9 & 19 AP Chapters ; 17

2. Oxidation-Reduction (“Redox”)
Ionic compounds are formed through the transfer of electrons
An oxidation-reduction rxn involves the transfer of electrons
We need a way of keeping track – oxidation states

3. Oxidation States A way of keeping track of the electrons.
Not necessarily true of what is in nature, but it works.
Use “+2” instead of “2+” since not necessarily actual charge
need the rules for assigning
memorize these!

4. Rules for assigning oxidation states
The oxidation state of an atom in an element is zero (i.e. Na(s), O2(g), O3(g), Hg(l))
Oxidation state for monoatomic ions are the same as their charge. (i.e. Na+, Cl-)
Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide (such as H2O2).
In compounds with nonmetals hydrogen is assigned the oxidation state +1.
In its compounds fluorine is always –1.
The sum of the oxidation states must be zero in compounds or equal the charge of the ion.

5. Oxidation States Practice Assign the oxidation states to each element in the following.
CO2
NO3-
H2SO4
Fe2O3
Fe3O4 +4 -2 +3 -2 +5 -2 -2
+8/3 -2 +1 +6 -2

6. Oxidation-Reduction
Electrons are transferred, so the oxidation states change.
2Na + Cl2 ® 2NaCl
CH4 + 2O2 ® CO2 + 2H2O

7. OIL RIG Oxidation-Reduction Oxidation is the loss of electrons.
Reduction is the gain of electrons.

8. Oxidation-Reduction LEO GER Losing electrons - oxidation
Gaining electrons - reduction

9. Oxidation-Reduction
Oxidation means an increase in oxidation state - lose electrons.
Reduction means a decrease in oxidation state - gain electrons.
The substance that is oxidized is called the reducing agent.
The substance that is reduced is called the oxidizing agent.

10. Redox Reactions

11. Agents Oxidizing agent gets reduced. Reducing agent gets oxidized.
Gains electrons.
More negative oxidation state.
Reducing agent gets oxidized.
Loses electrons.
More positive oxidation state.

12. Practice In the following reaction, identify the… Oxidizing agent; Reducing agent; Substance oxidized; Substance reduced
Fe (s) + O2(g) ® Fe2O3(s)
(+3) oxidation +3 -2
(-2) reduction
OA = O2
RA = Fe

13. Fe2O3(s)+ 3 CO(g) ® 2 Fe(l) + 3 CO2(g)
Practice In the following reaction, identify the… Oxidizing agent; Reducing agent; Substance oxidized; Substance reduced
Fe2O3(s)+ 3 CO(g) ® 2 Fe(l) + 3 CO2(g)
(+2) oxidation +3 -2 +2 -2 +4 -2
(-3) reduction
OA = Fe
RA = C

14. SO32- + H+ + MnO4- ® SO42- + H2O + Mn2+
Practice In the following reaction, identify the… Oxidizing agent; Reducing agent; Substance oxidized; Substance reduced
SO32- + H+ + MnO4- ® SO42- + H2O + Mn2+
(+2) oxidation +4 -2 +1 +7 -2 +6 -2 +1 -2 +2
(-5) reduction
OA = MnO4-
RA = SO32-

15. Half-Reactions
All redox reactions can be thought of as happening in two halves.
One produces electrons - oxidation half.
The other requires electrons - reduction half.

16. Half-Reactions Practice
Write the half reactions for the following:
Na + Cl2 ® Na+ + Cl-
(+1) oxidation +1 -1
(-1) reduction
Oxidation:
Na  Na+
Reduction:
Cl  Cl-
+ e- e- 2

17. Half-Reactions Practice
Write the half reactions for the following:
SO32- + H+ + MnO4- ® SO42- + H2O + Mn2+
(+2) oxidation +4 -2 +1 +7 -2 +6 -2 +1 -2 +2
(-5) reduction
Oxidation:
SO  SO42-
Reduction:
MnO  Mn2+
Clearly there’s more to this than just adding electrons…

18. Balancing Redox Equations
In aqueous solutions the key is the number of electrons produced must be the same as those required.

19. Acidic Solution For reactions in acidic solution an 8 step procedure:
Write separate half reactions
For each half rxn, balance all reactants except H and O
Balance O using H2O
Balance H using H+
Balance charge using e-
Multiply equations to make electrons equal
Add equations and cancel identical species
Check that charges and elements are balanced.

20. Redox Balancing Practice
Balance the following redox rxn that takes place in acidic sol’n:
(+6) oxidation
acid
Cr2O72-(aq) + C2H5OH(l) ® Cr3+(aq) + CO2(g) +6 -2 -2 +1 -2 +1 +3 +4 -2
(-3) reduction
Oxidation:
3H2O +
C2H5OH  CO2 2
+ 12H+
+ 12e-
Reduction: 2
6e- +
14H+ +
Cr2O72-  Cr3+ 2
+ 7H2O
12e- +
28H+ +
2Cr2O72-  Cr3+ 4
+ 14H2O

21. Redox Balancing Practice
Balance the following redox rxn that takes place in acidic sol’n:
Now add the oxidation and reduction reactions together:
3H2O +
C2H5OH  CO2 2
+ 12H+
+ 12e- 16 11
12e-+
28H+ + 2
Cr2O72- Cr3+ 4
+ 14H2O
16H+ + 2Cr2O72- + C2H5OH  2CO2 + 4Cr H2O

22. Practice
The following reactions occur in aqueous solution. Balance them
Cr(OH)3 + OCl- + OH- ® CrO4-2 + Cl- + H2O
MnO4- + Fe+2 ® Mn+2 + Fe+3
Cu + NO3- ® Cu+2 + NO(g)
Pb + PbO2 + SO4-2 ® PbSO4
Mn+2 + NaBiO3 ® Bi+3 + MnO4- + Na+

23. Now for a tough one…
Fe(CN)6-4 + MnO4- ® Mn+2 + Fe+3 + CO2 + NO3-

24. Basic Solution
Do everything you would with acid, but add one more step.
Add enough OH- to both sides to neutralize the H+ (OH- and H+ combine to form H2O)

25. Redox Balancing Practice
Balance the following redox rxn that takes place in basic sol’n:
(+1) oxidation
base
Ag(s) + CN-(aq) + O2(g) ® Ag(CN)2-(aq) +2 -3 +1 +2 -3
reduction?
Oxidation: 4 2
CN-(aq) + Ag(s)  Ag(CN)2-(aq)
+ e- 8
CN-(aq) + 4Ag(s)  4Ag(CN)2-(aq)
+ 4e-
Reduction:
4e- +
4H+(aq) +
O2(g) 
2H2O(l)

26. Redox Balancing Practice
Balance the following redox rxn that takes place in acidic sol’n:
Now add the oxidation and reduction reactions together: 8
CN-(aq) + 4Ag(s)  4Ag(CN)2-(aq)
+ 4e-
4e- +
4H+(aq) +
O2(g) 
2H2O(l)
8CN-(aq) + 4Ag(s) + 4H+(aq) + O2(g)  4Ag(CN)2-(aq) + 2H2O(l)
+4OH-(aq)
+4OH-(aq) 2
=4H2O(l)
8CN-(aq) + 4Ag(s) + 2H2O(aq) + O2(g)  4Ag(CN)2-(aq) + 4OH-(aq)

27. Redox Titrations Same as any other titration.
The permanganate ion is used often because it is its own indicator.
MnO4- is purple, Mn+2 is colorless. When reaction solution remains clear, MnO4- is gone.
Chromate ion is also useful, but color change (orangish-yellow to green) is harder to detect.